A significant detriment to the energy density of most batteries is posed by the battery's cathode. This is true for battery chemistries using, for example, lithium or nickel. Typically, oxidant is stored at the cathode at a molar charge capacity that is two to five times lower than that of the anode. Many fuel cells, on the other hand, use oxygen from the air as a source of oxidant. The existence of a continuous and virtually limitless oxidant source enables, in principle, high energy density. However, the use of hydrogen and organic fuels precludes high energy efficiencies due to problems with vapor pressure and balance-of-systems complexity, such as humidification and membrane issues. Metal-air electrochemical cells are able to combine the ultra-high anode capacity of batteries with the air-breathing cathode of fuel cells in order to achieve substantial energy densities that are relevant to modern energy demands.
Metal-air batteries typically include a fuel electrode at which metal fuel is oxidized, an air electrode at which oxygen is reduced, and an electrolyte solution for providing ion conductivity. A limiting factor with metal-air batteries is the evaporation of the electrolyte solution, particularly the evaporation of the bulk solvent, such as water in an aqueous electrolyte solution. Because the air electrode is required to be air permeable to absorb oxygen, it is also may permit the solvent vapor, such as water vapor, to escape from the cell. Over time, the cell becomes incapable of operating effectively because of this issue. Indeed, in many cell designs this evaporation issue renders the cell inoperable before the fuel is consumed. And this issue is exacerbated in secondary (i.e., rechargeable) cells, because the fuel may be re-charged repeatedly over the life of the cell, whereas the electrolyte solution is not (absent replenishment from an external source). Also, in rechargeable cells the water solvent is typically oxidized to evolve oxygen during re-charge, which may also deplete the solution.
There are other problems associated with conventional aqueous electrolyte batteries, such as water electrolysis during recharging, and self discharge. During recharge, a current is passed through the battery to reduce the oxidized fuel at the fuel electrode. Some of the current, however, electrolyzes the water resulting in hydrogen evolution (reduction) at the fuel electrode and oxygen evolution (oxidation) at the oxygen electrode as represented in the following equations:Reduction: 2H2O(l)+2e−→H2(g)+2OH−(aq) and  (1)Oxidation: 2H2O(l)→O2(g)+4H+(aq)+4e−  (2)In this manner, further aqueous electrolyte is lost from the battery. Additionally, the electrons that are consumed in reducing hydrogen are not available to reduce the fuel oxide. Therefore, the parasitic electrolysis of the aqueous electrolyte reduces the round trip efficiency of the secondary battery.
Self-discharge may result from impurities in the electrodes or reaction with the electrolyte. Typically, self-discharge from impurities in the electrodes is small (2-3% loss per month). The reaction of an active metal with water and/or O2 dissolved in the water, however, may be quite high (20-30% per month).
To compensate for these problems, metal-air batteries with aqueous electrolyte solutions are typically designed to contain a relatively high volume of electrolyte solution. Some cell designs even incorporate means for replenishing the electrolyte from an adjacent reservoir to maintain the electrolyte level. However, either approach adds significantly to both the overall size of the cell, as well as the weight of the cell, without enhancing the cell performance (except to ensure that there is a sufficient volume of electrolyte solution to offset evaporation of the water or other solvent over time). Specifically, the cell performance is generally determined by the fuel characteristics, the electrode characteristics, the electrolyte characteristics, and the amount of electrode surface area available for reactions to take place. But the volume of electrolyte solution in the cell generally does not have a significant beneficial effect on cell performance, and thus generally only detracts from cell performance in terms of volumetric and weight based ratios (power to volume or weight, and energy to volume or weight). Also, an excessive volume of electrolyte may create a higher amount of spacing between the electrodes, which may increase ionic resistance and detract from performance.
Another problem that arises with alkaline electrolyte cells is the formation of filaments or dendrites during the charging/discharging cycle. For example, during the charging of a rechargeable electrochemical cell, metal cations in the electrolyte are reduced at the electrode and are electrodeposited onto the electrode as the metal. Ideally, the electrodeposited metal is laid down as a smooth layer over the entire electrode surface, thereby preserving the electrode surface morphology from one discharge-charge cycle to the next. In practice, however, the metal tends to deposit preferentially at certain sites on the electrode. As a consequence, the morphology of the metal deposit is such that the electrode surface undergoes modification ranging from moderate roughening to formation of a coating of filaments or dendrites over the entire surface. After several cycles, the electrode can become covered by a dense mat of interwoven dendrites. This type of metal deposition is undesirable and also hazardous in electrochemical cells because the metal dendrites are often small enough to penetrate the microporous materials that are conventionally used to separate the anode from the cathode current collector. As a consequence, the dendrites can grow through the separator material and cause a short-circuit between the electrodes, resulting in cell failure and possible explosion. Dendrite growth around the edges of the separator material can also occur with similar results.
The use of non-aqueous systems for electrochemical cells has been suggested (see, e.g., U.S. Pat. No. 5,827,602). In non-aqueous systems, the aqueous electrolyte may be replaced with an ionic liquid. Ionic liquids which contain a strong Lewis acid such as AlCl3, however, are known to liberate toxic gases when exposed to moisture.
The use of low or room temperature ionic liquid rather than an aqueous electrolyte in a metal-air electrochemical cell, as described in U.S. Provisional Application Ser. No. 61/383,510, filed Sep. 16, 2010; 61/355,081, filed Jun. 15, 2010; 61/334,047, filed May 12, 2010; 61/329,278, filed Apr. 29, 2010; 61/177,072, filed May 11, 2009, and 61/267,240, filed Dec. 7, 2009, and described in U.S. patent application Ser. No. 13/105,794, filed on May 11, 2011; Ser. No. 13/096,851, filed Apr. 28, 2011; Ser. No. 13/085,714, filed Apr. 13, 2011; and Ser. No. 12/776,962, filed May 10, 2010, the disclosures of each of which are incorporated herein by reference in their entirety. The use of a low or room temperature ionic liquid in the cell essentially eliminates the problems associated with evaporation of solvent from an electrolytic solution.
Room temperature ionic liquids have extremely low vapor pressures (some have vapor pressures that are essentially immeasurable under standard conditions) and thus experience little or no evaporation. Therefore, cells using low or room temperature ionic liquids as their ionically conductive media need not incorporate excessive volumes of solution in order to compensate for evaporation over time. Relatively small amounts of ionic liquid are sufficient to support the electrochemical reactions needed for cell operation, thereby reducing cell weight and volume and increasing power to volume/weight ratios. Also, other problems associated with solvents, such as hydrogen evolution in an aqueous solution, may be avoided. This development is not conceded to be prior art and merely is described for contextual purposes to facilitate an understanding of the further development described herein.
Quaternization of tertiary amines, especially imidazole compounds and linear amines, by reaction with dimethyl carbonate is known. Some describe the use of the carboxylates or carbonates produced by this reaction as useful ionic liquid intermediates. See, e.g., Holbrey, et al., “1,3-Dimethylimidazolium-2-carboxylate: the unexpected synthesis of an ionic liquid precursor and carbine-CO2 adduct,” Chem. Commun. (2003), 28-29; Smiglak, et al., “Ionic Liquids via reaction of the zwitterionic 1,3,-dimethylimidazolium-2-carboxylate with protic acids, . . . ” Green Chem., Vol. 9, pp. 90-98 (2006); Bridges, et al., “An Intermediate for the Clean Synthesis of Ionic Liquids . . . ,” Chem. Eur. J., 5207-5212 (2007); Holbrey, et al., “Optimized microwave-assisted synthesis of methylcarbonate salts; . . . ,” Green Chem., Vol. 12, pp 407-413 (2010); Yang, et al., “Dimethyl Carbonate Synthesis catalyzed by DABCO-derived basic ionic liquids via transesterification of ethylene carbonate with methanol,” Tetrahedron Letters, 51, pp 2931-2934 (2010); and U.S. Pat. Nos. 4,892,944, and 5,865,513, the disclosures of each of the U.S. patents is incorporated by reference herein in its entirety.